Structure and Bonding
Year 12
Chemistry A.S. 2.4 (91164)
5 External credits
ATOMS Recap
Atoms are the building blocks of all matter.
An atom is made up of smaller bits called…
Name
Protons
Electrons
Neutrons
Charge
Positive
Position
Mass
negative
Nucleus
Outside
1
0
neutral
Nucleus
1
ELECTRON ARRANGEMENT
Electrons are very fast moving. They are arranged in
quantised energy levels / shells around the nucleus.
Electron shells determine how an atom behaves when it
encounters other atoms
The first shell fits…
2e
The second fits…
8e
The third fits…
8e
So the electron shell for
12Mg
would be… 2, 8, 2
Valence electrons increase across a period and remain
the same down a group.
Periods represent individual energy levels.
Gaining or losing electrons forms ions.
Octet Rule = atoms tend to gain, lose or share electrons so
as to have 8 in their valence shell and be stable.
C needs to
N would like to
O would like to
Share 4 electrons – why?
Gain 3 electrons
Gain 2 electrons
Lewis Diagrams
Symbols of atoms with dots to represent the
valence-shell electrons (electron dot structures)
1
2
13
14
15
16
17
H
Li
18
He:


Be
B 


C


Na Mg


Al 


N




O



 Si 
P 
S 





: F  : Ne:




:Cl  :Ar :


Learning Check

A.
X would be the electron dot formula for
1) Na
B.

 X 

1) B
2) K
3) Al
would be the electron dot formula
2) N
3) P
Sum of
valence e-
:
For NF3
: F:
: F:
:
Atom
placement
N
Lewis
structure
N
5e-
F 7e- X 3 = 21eTotal
26e-
:
: F:
Remaining
valence e-
Place atom with
most incomplete
shell in the centre
:
Molecular
formula
Draw single bonds.
Subtract 2e- for each bond.
Give each atom 8e(2e- for H)
Creating a Lewis diagram
Lewis Diagrams for Molecules with
Multiple Bonds
Write Lewis structures for the following:
(a) Ethylene (C2H4)
(b) Nitrogen (N2)
If a central atom does not have a full octet, then two e(either single or non-bonded pair) can be moved in to form
a multiple bond. B (6) and Be (4) are exceptions
(a) There are 2(4) + 4(1) = 12 valence e-. H can have only
one bond per atom.
SOLUTION:
H
H
:
H
C
C
H
H
Non-bonded e- become a double bond
C
H
H
C
(b) N2 has 2(5) = 10 valence e-. Therefore a triple bond is required to make
the octet around each N.
N
.
:
N
.
:
:
:.
N
.
N
N
:
.:
N
.
H
Electronegativity Values
• Ability of an atom to attract a pair of
bonded electrons
• EN range 0.7 (weak pull) - 4 (strong pull)
• EN increases across a period
• EN decreases down a group
• Group 18 atoms have EN value 0 – Why?
• Remember F,O,N,Cl
Predicting the type of Bonding
• Small or no difference in EN = covalent
• Large difference (>2) in EN = ionic
EN and the Periodic Table
Determining Bond
Polarity
•
•
•
•
Non-Polar
Bonding electrons are
shared equally
No difference in EN
Usually seen when
identical atoms bond
E.g. H2
•
•
•
•
Polar
Bonding electrons are
shared but not equally
Small difference in EN
(0.5-2)
Usually seen when 2
different non-metals bond
E.g. HCl
NON-POLAR COVALENT BONDS
• when electrons are shared
equally e.g.
H2 or Cl2
• EN of both atoms is the same
Non-Polar covalent bonds
Two atoms share one or more pairs of outer-shell electrons.
Oxygen Atom
Oxygen Atom
Oxygen Molecule (O2)
POLAR COVALENT BONDS
• when electrons are unequally
shared e.g.
HBr or H2O
• There is a small EN difference
between the atoms.
Polar covalent bonds
- water has polar bonds because… oxygen is more
electronegative than hydrogen, and therefore electrons
are pulled closer to oxygen forming a bond dipole.
Molecular Shape (VESPR)
• Electron density clouds repel each other to
give maximum separation and this
determines the molecular shape.
• 4 areas of e- density = Tetrahedral shape
• 3 areas of e- density = trigonal planar
shape
• 2 areas of e- density = linear shape
Molecular Shapes
Electron density cloud repulsions around the central atom
form three basic molecular shapes.
linear
trigonal planar
tetrahedral
Molecule shape with two electron clouds
Examples:
Cl2, CO2,
HCN
Molecule shapes with three electron clouds
Non bonding
electrons
AX3
Bond angle?
Examples:
Examples:
SO3, BF3, NO3-,
CO32-
SO2, O3, PbCl2,
SnBr2
Molecule shapes with four electron clouds
Maximum
repulsion
between
areas of
electron
density
NH3
PF3
ClO3
What if
there was a
non bonded
pair of
electrons?
H3O+ 2 non bonded
pairs?
Examples:
CH4, SiCl4,
SO42-, ClO4-
H2O
OF2
SCl2
Bond angle?
Exam Question
• Compare the molecular shape of water with carbon
dioxide. In your answer consider:
• Areas of electron density
• Bonding and non bonding electrons
• Bond angles
• Water has 4 areas of electron density around
the central atom which repel maximally. 2 areas
are bonded and 2 non bonded which results in a
bent shape with bond angle 109.5 whereas,
Carbon dioxide has 2 areas of electron density
around the central atom which repel maximally.
Both areas are bonded which results in a linear
structure with bond angle 180.
Molecular Polarity
• Molecular polarity occurs if the individual
bond polarities give the molecule a slight
charge (dipole).
-water is a polar molecule
because…
oxygen is more electronegative
than hydrogen, and therefore
bond electrons are pulled closer
to oxygen. The shape of the
molecule means the bond polarity
is not symmetrical, so dipoles
will not cancel.
Predicting the Polarity of Molecules
Draw the shape, find the EN values, determine the polarity of the
bonds and whether the molecule is polar or not.
molecule is
asymmetrical
so bond dipoles
don’t cancel
and molecule is
polar
(a) Ammonia, NH3
The shape is trigonal pyramidal.
H
N
H
H
H
N
H
H
ENN = 3.0
ENH = 2.1
bond polarity
H
N
H
molecular
polarity
H
Predicting the Polarity of Molecules
(b) Boron trifluoride, BF3
has 24 valence e- and all electrons around the B will be
involved in bonds.
The shape is trigonal planar.
Bond angle? Bond polarity?
ENF = 4.0
ENB = 2.0
F
B
1200
F
F
B and F have different EN, hence the
bonds are polar. The shape is
symmetrical so the dipoles cancel and
the molecule is non polar
Intermolecular bonding
• Intra-molecular polar
covalent bonds are
strong within the water
molecule.
• The individual
molecules are attracted
to each other by weak
inter-molecular forces.
• Breaking intermolecular bonds results
in phase changes.
Ionic bonds
Covalent bonds
Metallic bonds
Ionic Bond
• Bond formed
between
two ions by
the transfer of
electrons
Ionic Bond
• Between atoms of metals and non-metals
• Due to very different electronegativities
• Bond is formed by the transfer of electrons
and is very strong
• Produce charged ions all states.
• Good conductors
• Have high melting point.
• Examples; NaCl, CaCl2, K2O
Cations - Ions from Metals
 Ionic compounds result when metals react with
nonmetals
 Metals lose electrons to form a complete
valence shell.
 Positive ions form when the number of
electrons are less than the number of protons
Group 1 metals 
ion 1+
Group 2 metals 
ion 2+
Group 13 metals 
ion 3+
Formation of Sodium Ion
Sodium atom
Na 
2-8-1
11 p+
11 e0
– e
Sodium ion

Na +
2-8 ( = Ne)
11 p+
10 e1+
Some Typical Ions with
Positive Charges (Cations)
Group 1
Group 2
Group 13
H+
Mg2+
Al3+
Li+
Ca2+
Na+
Sr2+
K+
Ba2+
Learning Check
A. Number of valence electrons in aluminum
1) 1 e2) 2 e3) 3 eB.
Change in electrons for octet
1) lose 3e2) gain 3 e3) gain 5e-
C.
Ionic charge of aluminum
1) 32) 5-
3) 3+
Anions - Ions from Non-metals
In ionic compounds, non-metals in group
15, 16, and 17 gain electrons from metals
Non-metal atoms add electrons to achieve a
full octet
Non-metal ionic charge:
3-, 2-, or 1-
Formation of Fluoride Ion
Fluoride atom

:F

+ e
Fluoride ion
 1-
: F:

2-7
2-8 (= Ne)
9 p+
9 e0
9 p+
10 e1-
1). Ionic bond – electron from Na is transferred to Cl,
this causes a charge imbalance in each atom. The Na
becomes (Na+) and the Cl becomes (Cl-), charged
particles or ions.
Sodium Chloride Ionic Bond
Ionic solid
• 3D Lattice of positive
(cations) and negative
(anions)
• Held together by
strong electrostatic
attraction between the
ions called ionic
bonding
• E.g. Salt
N.B. Name the ions in
exams
Properties of Ionic Solids
• Melting Point
– High (lots of energy
needed to break
strong bonds)
• Solubility
– Most dissolve in water
(strong attraction of
polar water molecules
overcomes the ionic
bond)
• Conductivity
– Do not conduct electricity
(ions are not free to
move)
– If dissolved or molten they
will conduct
• Hardness
– Hard but brittle (a force
can cause layers of ions
to shift so they repel each
other)
Definitions
• Electricity – movement of charged
particles
• Melting – If the energy supplied
overcomes the attractive force between
the particles it will melt
• Dissolving – If the solvent/solute attraction
is greater than the solvent/solvent and
solute/solute attractions then the
substance will dissolve
METALLIC BOND
• bond found in
metals
• holds metal
atoms together
very strongly
Metallic Bond
• Formed between atoms of metallic
elements
• Free flowing “sea of non directional
electrons” around the metallic cations
• Good conductors at all states
• Lustrous
• Very high melting points
• Non-soluble
• Examples; Na, Fe, Al, Au, Co
Metallic Bond
A non-directional sea of electrons
Metallic Solid
• 3D lattice of metallic
atoms surrounded by
a sea of valence
electrons
• Held together by the
attraction of the metal
cations for the
valence electrons
(metallic bonding)
• E.g. Gold
Properties of Metallic Solids
• Melting Point
– Varies, but often high
(lots of energy needed
to break metallic
bonds)
• Solubility
– Not soluble. Strong
bonds are not broken
by the attraction of
solvent molecules
• Conductivity
– Very good (due to the free
moving electrons)
• Hardness
– Can be deformed without
breaking (due to the nondirectional bonds)
– Metals can be hammered
into shape (malleable)
and drawn into wires
(ductile)
COVALENT BOND
• Bond formed
by the
sharing of
electrons to
achieve a
full octet.
Covalent Bond
• Between non-metallic elements of similar
electronegativity (EN).
• Formed by sharing electron pairs
• Stable non-ion forming particles, they are
not conductors at any state
• May be polar or non-polar depending on
EN difference
• Examples; O2, CO2, C2H6, H2O, SiC, HCl
Covalent Network Solids
• Linear chains, 3D
lattice or 2D
layers of atoms
• Held together by
extremely strong
covalent bonds.
• E.g. Diamond
and Graphite
Properties of Covalent Network Solids
• Melting Point
– Extremely high (lots of
energy needed to
break strong covalent
bonds)
• Solubility
– Not soluble. Strong
bonds are not broken
by the attraction of
solvent molecules
• Conductivity
– Do not conduct electricity
(no free moving charged
particles)
– Except Graphite (due to
the movement of free
electrons between the 2D
planes)
• Hardness
– Extremely hard
– Graphite is brittle
Covalent Network Solids
Linear Chains
• Polymers such as
polypropylene
• Larger molecular
mass means more
WiFs between chains,
therefore M.P.
increases
• No charged particles
means low electrical
conductivity
• Chains can slide over
one another so
substance is soft and
flexible
Molecular Solid
• Formed from
individual
covalently bonded
molecules
• Held together by
weak
intermolecular
forces
• E.g. ice, solid CO2
Properties of Molecular Solids
• Melting Point
– Low (little energy
needed to break weak
intermolecular forces)
– Higher in heavy or
more polar molecules
(Wif forces are
stronger)
• Solubility
– Depends on polarity
(like dissolves like)
• Conductivity
– Do not conduct electricity
(there are no charged
particles)
• Hardness
– Easily broken
(intermolecular forces are
weak)
Trends in the Periodic Table
Numbers represent atomic radii
Chloride compounds of period 3
elements
• NaCl MgCl2 AlCl3 SiCl4 PCl3 SCl2 Cl2
Ionic compound is
formed by the
transfer of
electrons due to
high
electronegativity
differences
(metal/non-metal)
Molecular
compound is
formed by the
sharing of
electrons due to
lower
elecronegativity
differences (2x
non-metal)
Trends down group 17 (Halogens)
Name
Fluorine
Atomic M.P. State (room
0C
size
temp)
0.64
-222 Yellow gas
Chlorine 0.99
Bromine 1.14
-101 Yellow green
gas
-7
Orange liquid
Iodine
1.33
114
Astatine
1.4
302 radioactive
Grey solid
Weight
heavier
Energy Changes
• CH4 + 2O2
CO2 + 2H20 ∆H = -888kJmol-1
• The intramolecular bonds between C and H atoms in
Methane and between O atoms in Oxygen must
break.
• Breaking bonds requires energy
• The free atoms now rearrange and new
intramolecular bonds are made between C and O
atoms forming Carbon Dioxide and H and O atoms
forming water.
• Making bonds releases energy
• The difference between bond making and bond
breaking means overall this reaction releases energy
Enthalpy ∆H
• Enthalpy is the energy in a substance
enthalpy change
during a reaction
can be measured
as energy is
released or
absorbed from
the surroundings
• ∆H is the difference between the energy needed
to break the bonds in the reactants, and the
energy given out when new bonds are made in
the products. It is measured in Kj/mol.
Enthalpy of
products and
reactants
cannot be
measured,
but…
∆H reaction = ∑E bonds broken - ∑E bonds made
• Bond breaking is
always endothermic
• The energy required
to break bonds
(dissociation enthalpy)
depends on the
strength of the bond.
• Double bonds are held
tighter than single
bonds so the atoms
are closer together.
• Bond energies can be
calculated from the
heat of reaction.
Bond Energies
Read ESA pg 145
complete activity 11c
Exothermic Reactions
Exothermic Reactions
Energy is given out (exits) from
the chemical bonds
The products have less
energy than the
reactants
Surroundings gain heat energy
Enthalpy Change is negative (∆H-)
Combustion and neutralisation reactions are
exothermic
Endothermic Reactions
Endothermic Reactions
Energy is taken in (enters) the
chemical bonds
The products have
more energy than
the reactants
Surroundings lose heat energy
Enthalpy Change is positive (∆H+)
Photosynthesis or gas
liquid
solid
Surroundings
Exothermic reaction ∆HEnergy
Energy
Bonds made
Bonds made
Solid
Liquid
Bonds broken
Energy
Gas
Bonds broken
Energy
Endothermic reaction ∆H+
Summary Table
Exothermic
reactions
Endothermic
reactions
Energy is given out to
the surroundings
Energy is taken in from
the surroundings
∆H is negative
∆H is positive
Products have less
energy than reactants
Products have more
energy than reactants
Calculating ∆E
• Energy is measured in kJ or joules
• CH4 + 2O2
CO2 + 2H20 ∆H = -888kJmol-1
Enthalpy change for this reaction i.e. 888kJ of
heat is released when 1 mol methane reacts
with 2 mols of oxygen to produce 1 mol carbon
dioxide and 2 mols of water
How much heat is released if 1.5mols of CH4 burns?
1332kJ
Calculating ∆E
• CH4 + 2O2
CO2 + 2H20 ∆H = -888kJmol-1
• How much energy is released when 1kg of
CO2 is produced?
M( C ) = 12gmol-1, M (O ) = 16gmol-1
m = 1000g, M (CO2) = 44gmol-1
Moles of CO2 produced n=m/M
n= 1000/44
n= 22.72 moles
888kJ are produced to make 1 mole CO2 therefore 888
x 22.72 =20181.81kJ
Calculating amount (moles)
CH4 + 2O2
CO2 + 2H20 ∆H = -888kJmol-1
• How many moles of water are produced when
10000kJ of energy are released?
2 moles produced when 888kJ are
released
Therefore 1 mole produced for
every 444kJ
10000/444 = 22.5moles
Calculate ∆H
• 0.01g Mg(s) reacts with HCl and the
temperature rises producing 100.8J.
Calculate the enthalpy of reaction in kJmol-1
Moles of Mg
n=m/M
n= 0.01/24
n= 4.167 x 10-4 moles
• ∆H = ∆E /n
= 100.8/4.167 x 10-4
= -242kJmol-1